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Chemistry > semester 1 exammm > Flashcards

semester 1 exammm Flashcards

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1
Q

what is atomic radius

A

distance from outermost electron to nucleus

atomic radius decreases left to right across a period due to:

atom’s increasing nuclear charge &

high nuclear charge increases attraction of electrons, bringing them closer to the nucleus therefore decreasing atomic radius

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2
Q

what is ionisation energy

A

amount of energy needed to remove an electron from a neutral atom in the gaseous state

e.g. metals like lithium have low ionisation energy

non metals have high ionisation energy as they almost have a full valence shell

ionisation energy increases from left to right across a period

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3
Q

what are the 3 factors that explain the trends of the periodic table

A
  1. atom’s nuclear charge: great nuclear charge the more strongly electrons are attracted to the nucleus
  2. atomic radius: as the distance between the nucleus and outermost electron increases, the strength of electron attraction decreases
  3. sheilding by inner shell electrons: electrons in the inner shell repel other electrons between the nucleus and the outer electrons, blocking some of the atom’s electronegativity to the outer electrons
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4
Q

what is effective nuclear charge

A

number of protons in nucleus - number of inner shell electrons

increases left to right: as nucleus becomes more positive attraction between nucleus and electrons becomes stronger (increasing effective nuclear charge.)

more energy needed

decreases down group: electrons get further from nucleus and are less strongly bound - sheilding effect is greater

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5
Q

what is successive ionisation energy

A

energy needed to remove subsequent electrons

values become larger because electrons are being removed from ions which have increasing positive charges

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6
Q

explain electronegativity

A

definition: the electron attracting power of an atom in a molecule

metals have low electronegativity values

non metals have high electronegativity values

increases across a period:

decreases down a group:

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7
Q

what is emission spectra

A

when atoms absorb energy e.g. heat they become exited & can move up to a higher energy level (an e- shell thats further away)

when the e- moves back into its original position it releases light energy in all different directions

the light will have a specific wavelength unique to each element

for each particular element, all atom energy levels are identical

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8
Q

what are the 3 different types of spectra

A

emission spectroscopy

atomic absorption spectroscopy (AAS)

mass spectrometry

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9
Q

explain what atomic absorption spectroscopy (AAS) is

A

used to perform quantitative analysis (find amount of element present)

relies on electron transfer between atomic energy levels

uses the absorption of light by the atoms electrons to measure how much of an element is present

the lamp used for this process is made out of the same element being tested

e.g. if you were testing for zinc, you would use a zinc lamp

the light emitted has a unique set of wavelengths unique to the element.

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10
Q

what is the best analytical technique for determining the isotopic concentration of a sample

A

atomic absorption spectroscopy

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11
Q

explain the process of an atomic absorption spectrometer

A
  1. an electric current is passed through a gaseous sample of the element so it will emit light
  2. the light emitted has a unique set of wavelengths
  3. the sample is vaporised
  4. the light from the lamp is passed through the vaporised sample
  5. the element being tested for will absorb the light
  6. light enters the monochromator which focuses on 1 wavelength by the detector
  7. detector measures amount of light passed through the sample that is not absorbed
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12
Q

what is an isotope

A

atoms that have the same atomic number but different mass number (number of protons)

share the same chemical properties (due to same number of valence electrons)

physical properties (mass) will differ

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13
Q

what is relative atomic mass

A

actual atomic mass is too small to calculate

it is the mass of the atom compared to 1 twelfth of the mass of a carbon-12 atom

its an average of all naturally occurring isotopes present

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14
Q

list the properties of metal

A

good conductors of heat and electricity in solid & liquid state

shiny surface

hard

malleable

ductile

insoluble in water

reasonably high MP

very high BP

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15
Q

describe metallic bonding

A

particles arranged as a lattice of positive cations surrounded by a sea of delocalised electrons

delocalised: can move around atoms and through the lattice

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16
Q

explain the metallic property of heat/electricity conduction

A

when heated energy of cations and delocalised electrons increases

this energy is conducted along metals as cations and electrons collide, transferring electricity

when metals conduct electricity, electrons from one end of the metal cause a similar number of e- to be displaced from the other end

in sold: only electrons move
liquid: both cation and e- carry charge

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17
Q

explain the metallic property of malleability and ductility

A

bonding in metals is NON DIRECTIONAL

layers of atoms can slide over each other without destroying metallic bond

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18
Q

explain the metallic property of MP & BP

A

MP and BP indicate strength of intermolecular forces

relatively high MP and BP due to strong attraction between positive cations and mobile electrons

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19
Q

explain the metallic property of hardness

A

strength of bond is related to charge of positive cations in lattice

metals in group 2 are stronger because atoms release 2 valence electrons (instead of 1 in group 1) and form ions with a +2 charge

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20
Q

list the ionic properties

A

non conductors of electricity in solid state

good electrical conductors in molten or liquid state

hard

brittle

relatively high MP

relatively high BP

varying solubilities in water, insoluble in non-polar solvents

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21
Q

describe ionic bonding

A

ionic solid:
positive and negative ion are arranged in an orderly 3D lattice

each positive ion is surrounded by negative ions and vice versa

ions are held in fixed positions by strong ionic bonds

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22
Q

Explain the ionic property of conducting heat/electricity

A

solid: non conductors because ions are held in a fixed lattice, no mobile charged particles
molten: ions are now mobile, both positive and negative ions carry charge
aqueous: ions dissociate and are able to carry charge

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23
Q

Explain the ionic property of MP and BP

A

high MP and BP:

ionic bonding has large electrostatic forces of attraction between holding atoms together in a lattice

large amount of energy is required to overcome strong ionic bonds

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24
Q

Explain the ionic property of hardness

A

strong ionic bonding extends throughout lattice solid

scratching requires overcoming strong ionic bonds

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25
Q

Explain the ionic property of brittleness

A

ionic lattice doesn’t withstand distortion

when lattice is distorted, like charges align and repel, shattering lattice

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26
Q

Describe covalent molecular substances

A

exist only as discrete molecules

non metals only

solid, liquid or gaseous state

intramolecular forces: strong covalent bond hold atoms together within a molecule

intermolecular forces: weak forces of attraction between molecules

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27
Q

list the properties of CMS

A

non conductors of electricity in solid and liquid

some are weak/good conductors in aqueous solutions

malleable

soft

low Mp and BP

mostly insoluble in water, soluble in non-polar compounds

examples: F2, H2O, O2

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28
Q

explain the covalent molecular property of electricity conduction

A

non conductors in aq, s or l:

molecules are uncharged

electrons are held tightly within covalent bonds

**exception: some CMS react with water to form ions, and become conductors in aqueous solutions

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29
Q

explain the covalent molecular property of softness

A

weak intermolecular forces require little effort to overcome: easy to separate molecules

strong covalent bonds aren’t disrupted

30
Q

explain the covalent molecular property of malleability

A

molecules can easily slide or be pushed

only weak intermolecular forced exist between them, these can easily be reformed

31
Q

explain the covalent molecular property of BP and MP

A

low BP and MP:

when CMS are heated, little energy is needed to overcome weak intermolecular forces

32
Q

Describe covalent network substances

A

solid state only

non metals only

carbon with different structures are called allotropes:
diamond and graphite
examples of CNS substances: silicon, silicon dioxide, carbon

33
Q

List the properties of covalent network substances

A

non conductors of heat and electricity in solid and liquid state **exception: graphite

extremely brittle

extremely hard

high MP and BP

insoluble in water and most other solvents

reasonably chemically inert

34
Q

Describe the bonding of covalent network substances

A

arranges in an infinite lattice structure

35
Q

explain the CNS property of electricity/heat conduction

A

non conductors: no mobile electrons because all valence electrons are tied up in covalent bonds **exception: graphite

36
Q

explain the CNS property of brittleness

A

if some bonds are broken, the rest of the lattice is placed under stress and shatters

37
Q

explain the CNS property of hardness

A

extremely hard:

breaking strong covalent bonds requires a very large amount of energy

38
Q

explain the CNS property of MP and BP

A

high MP and BP:

each atom is held in place by strong covalent bonds

breaking these bonds requires a very large amount of energy

39
Q

explain the CNS property of solubility

A

insoluble in water & most other solvents:

vary large amount of energy is required to re-arrange existing bonds in any dissolving process

40
Q

explain the CNS property of chemical inertness

A

reasonably chemically inert:

all bonding requirement are fulfilled

very large amounts of energy required to re-arrange existing bonds

41
Q

Briefly describe the structure and property of diamond

A

structure:
each carbon atoms is attached to 4 other atoms in a 3D arrangement called a tetrahedron

properties:
high MP (>3550˚C)
non conductor of electricity
brittle
extremely hard
42
Q

Briefly describe the structure and property of graphite

A

structure:
each carbon atoms is covalently bonded to 3 other atoms to form flat 2D sheets called graphene

one valence electrons from each carbon remain delocalised

properties:
high MP
good conductor of electricity
soft

43
Q

what does the mole represent in mole/mass calculations

A

measure of the number of particles in a substance

1 mole of any substance contains 6.022 x 10(23) particles

44
Q

what does molar mass represent in mole/mass calculations

A

equal to atomic, molecule and formula masses in grams

45
Q

What is mass spectrometry used for

A

determines mass and relative abundance of each isotope present in a sample of an element

46
Q

describe the process of mass spectrometry

A 
involved 4 steps:
vaporisation
ionisation
acceleration
detection
  1. sample is vaporised
  2. electrons are fired at the sample and are used to knock off electrons from the atoms, forming ions
  3. ions are accelerated and are deflected through a magnetic field (heavier ions are deflected less than lighter ions)
  4. Ions are detected and produces a mass spectrum
47
Q

list 3 separation techniques

A

sieving

filtration

evaporation

48
Q

Briefly describe the separation technique of filtration

A

liquid passes through filter paper (filtrate) ad solid residue is left behind

used for:
separating very fine solid particles from liquids and gases

49
Q

Briefly describe the separation technique of sieving

A

mixture is passed through a sieve, a series of sieves may be used

based on different sizes of components, smaller particles will pass through the hole, larger particles will get caught

used for:
separating solids from liquids
e.g. straining pasta

50
Q

what are the classifications of matter

A

pure substances: contain 1 type of particle

  • elements: contain 1 type of atom e.g. Ca, O2
  • compounds: contain 2 or more different types of atoms
    e. g. H2O, NaCl

mixtures: contain 2 or more different types of particles

  • homogenous
    • uniform composition
    • e.g. air, soft drinks
  • heterogenous
    • non uniform composition
    • e.g. dirty water, paints
51
Q

what determines the shape of molecules

A

the arrangement of electrons pairs surrounding the central atom

52
Q

what does VSEPR stand for and what does this theory state

A

stands for valence shell electron pair repulsion

theory states that electrons pairs in the valence shell of an atom repel each other, and are arranges as far apart as possible

53
Q

what does polarity refer to

A

related to the sharing of the bonding electrons between atoms

54
Q

what is the difference between a polar and non polar bond

A

non polar: electrons are shared evenly between atoms, equal charge distribution

polar: electrons are not shared evenly between atoms because different atoms have different electronegativities

55
Q

explain what a polar bond is

A

occurs between atoms with 2 different electronegativities

electrons spend more time moving around the atom with the greatest electronegativity

this leads to an uneven charge distribution in the bond

the region near the more electronegative atoms has a partial negative bond

the region near the less electronegative atoms has a partial positive bond

this is called a bond dipole

56
Q

what is an extreme polar bond

A

occurs between non metal atoms with large differences in electronegativities e.g. F and H

57
Q

what are intermolecular forces

A

weak attractive forces existing between molecules

58
Q

list the different types of intermolecular forces

A

dispersion: attraction between all types of particles resulting in instantaneous dipoles

dipole-dipole: attraction between polar molecules

hydrogen bonds: attraction between polar molecules containing OH, NH or FH groups

ion-dipole: attraction between polar solvent water molecules and dissolved ions

59
Q

explain what dispersion forces are

A

temporary dipoles form between molecules in order for them to change shape

temporary polarity: electrons are constantly moving around atoms and are not always symmetrically distributed across the atom

this temporary asymmetrical distribution produces a temporary/instantaneous dipole

if it is close to another molecule w/ a temporary dipole, it will cause polarity in the neighbouring atom

these two temporary dipoles will be attracted to one another

this force between molecules is dispersion force

60
Q

how can you predict the strength of dispersion forces

A

increase in strength with increasing number of electrons
- as ayomic radius increases, e’ have greater space to
mover and greater chance to arrange asymmetrically

molecular shape

the larger the surface area over which the electrons can develop an instantaneous dipole and the closer molecules can fit next to each other, the strogner the dispersion force

61
Q

explain what dipole-dipole forces are

A

weak attractive forces between polar molecules

molecule that has a partial negative charge is attracted to neighbouring molecule with partial positive charge

the more polar the molecules, the stronger the force will be

62
Q

explain what hydrogen bonding is

A

dipole-dipole forces containing NH, OH or FH as they’re the most electronegative atoms

very polar bond –> stronger dipole dipole forces

more energy required to overcome bonds: higher BP

atomic radii are smaller - high charge density at the end of each molecule

e- are attracted more strongly toward N, O or F

63
Q

what are the requirements for hydrogen bonding

A

presence in 1 molecules of at least 1 NH OH or FH covalent bond

presence of another polar molecule containing at least 1 N, O or F atoms with a non-bonding pair of electrons

64
Q

define hydrogen bond

A

the attractive force between the lone pair of electrons on the N, O of F atom on on of the polar molecules and the partially positive charge of the H bonded to N, O of F on the second polar molecule

65
Q

what are the unique properties of water

A

relatively high MP (O˚) and BP (100˚)

density of water in the solid phase is lower than its density in liquid phase

high surface tension

66
Q

explain water’s property of its density

A

water has a maximum density of 1.0g ML-1 as a liquid at 4˚

when freezed it expands significantly and density drops by 8%

reduced density is why ice floats on water and freezes top down providing a layer of insulation preventing the rest of the water from freezing

this is because of the crystal lattice formed when water freezed - each water molecule can form 4 hydrogen bonds, maximising hydrogen bonding but takes up more space

67
Q

explain water’s property of its high surface tension

A

water has strong intermolecular forces

imbalance of these forces at waters surface causing water molecules to be pulled inward toward the bulk of the liquid

as a result the surface tries to contract to minimise surface area

68
Q

what is chromatography

A

separates compounds in a solution with the same properties

paper chromatography: separates a mixture of solutes with difference solubilities of absorption

69
Q

explain the process of chromatography

A

uses paper and a solvent that can move over the medium

samples are placed along chromatography paper and paper is dipped in a solvent and left to dry

solvent will run up the paper, dissolving the samples

different solubilities cause some samples to run further up the paper than others

liquid with the lowers BP is first to evaporate and be distilled out of the paper and collected

70
Q

what are the physical properties of intermolecular forces

A

changing phase

solubility

equilibrium vapour pressure

71
Q

explain the physical property of changing phase for intermolecular forces

A

MP and BP of substances increases as intermolecular force strength increases

as molecular mass increases (number of e’), MP and BP increases

substances that have dipole-dipole forces have a slightly higher MP and BP

substances that have hydrogen bonding have a slightly higher MP and BP

72
Q

explain the physical property of solubility for intermolecular forces

A

a solute and solvent can only form a solution if the attractive forces between molecules are comparable in strength

e,g, non polar molecules (dispersion forces) can’t be soluble in polar substances (dipole-dipole)

Chemistry (8 decks)

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